Study of complex formation between C₁₈H₃₆N₂O₆ and ${\text{UO}}_2^{2+}$ cation in some binary mixed non-aqueous solutions

1 Introduction

Uranium is a radioactive metal that exists ubiquitously in the environment (Gongalsky Konstantin, 2003). Since uranium is one of the main sources in nuclear energy generation and enriched uranium is a major component in nuclear weapons, human beings have a high chance of being exposed to uranium, which can cause severe effects on human health (Craft et al., 2004; Zhou and Gu, 2005). A large amount of uranium in drinking water may lead to harmful biological effects in humans. Chemical toxicity caused by natural uranium poses a major risk to the kidneys. Also it is important to monitor the concentration of uranium at various stages of preparation of uranium dioxide fuel pellets (Murty et al., 1997) and also in the environment safety assessment related to the nuclear industry (Miller et al., 1994). For these reasons, detection of uranium is very important. In nature, hexavalent uranium is commonly associated with oxygen as the uranyl ion, ${\text{UO}}_2^{2+}$ .

Thermodynamic studies on the complexation of C₁₈H₃₆N₂O₆ with the ${\text{UO}}_2^{2+}$ cation have been performed in order to elucidate the nature of the cation-binding behavior from the thermodynamic point of view and also to gain insights into the factors governing cation–ligand complexation phenomena (Davidson et al., 1984; Matthes et al., 1987; Arnold et al., 1987; Adamic et al., 1985, 1986). The ability of the solvent molecules to compete with the donor atoms of the ligand toward the cation, is one of the critical factors that can thermodynamically influence the complexation process (Solov'ev et al., 1996).

In order to achieve appropriate solvent properties, mixtures of two solvents are often used. The physiochemical properties of mixed solvents are interesting both from theoretical and practical points of view, because many chemical and electrochemical reactions might be carried out advantageously in these media. Usually mixed solvents do not behave as expected from statistical considerations; the solvating ability of solvents in their mixtures can be different from those of neat media. The deviations from ideal behavior are indicative of the extent of preferential solvation and the existence of specific solvent–solute and solvent–solvent interactions (Kinisinger et al., 1973; Burger, 1983).

Ion–solvent interactions play a very important role in the studies of stoichiometry, structure, and the stability of metal cation complexes with crown ethers. Knowledge of the ionophore solvation properties enables one to choose a suitable solvent for complexation studies and to obtain detailed information on the solvent effect. To control the chemical reactions in non-aqueous and mixed solvents, it is necessary to select an appropriate solvent (Marcus, 1985). This is a complicated problem because of an exceptionally great variety and specificity of chemical reactions in solutions. Recently, we have investigated the complexation processes of crown ethers with various metal cations in pure and some binary mixed organic solvent solutions (Rounaghi and Mohammad Zadeh Kakhki, 2009; Rounaghi et al., 2008, 2001, 2011; Razghandi et al., 2011; Shamsipur and Khayatian, 2001).

Uranium is at the heart of commercial nuclear power and is vital to the entire nuclear enterprise. Its presence and use worldwide has resulted in a vast literature, not only on uranium complexation and separation reactions, but also on all aspects of the uranium fuel cycle, mining and, milling, isotope enrichment, reactor fuel manufacturing and spent reactor fuel. The uranyl ion ${\text{UO}}_2^{2+}$ is the most common uranium ion in aqueous media. Uranyl ion complexes extract readily into organic solvents. This is the basis of the widely used purex process for reprocessing spent nuclear reactor fuel reprocessing and weapons production.

In the present work, we studied the complexation process between C₁₈H₃₆N₂O₆ (Fig. 1) and the ${\text{UO}}_2^{2+}$ cation in pure acetonitrile, tetrahydrofuran, dichloromethane and dimethylformamide and also in acetonitrile–tetrahydrofuran (AN–THF), acetonitrile–dichloromethane (AN–DCM) binary mixtures at different temperatures using the conductometric method.

Close

Figure 1

Structure of kryptofix 222.

Export to PPT

2 Experimental

C₁₈H₃₆N₂O₆ (Merck) and UO₂(NO₃)₂·6H₂O (Riedel company) were used without further purification. The solvents: acetonitrile, dimethylformamide, dichloromethane, tetrahydrofuran, all from (Merck), were used with the highest purity.

The experimental procedure to obtain the formation constant of the complex was as follows: a solution of metal salt (1 × 10⁻⁴ M) was placed in a titration cell and the conductance of the solution was measured, then a step-by-step increase in the C₁₈H₃₆N₂O₆ concentration was performed by a rapid transfer of C₁₈H₃₆N₂O₆ solution in the same solvent (2 × 10⁻³ M) to the titration cell, using a microburette and the conductance of the resulting solution was measured after each step at the desired temperature. The conductance measurements were performed on a conductivity meter LF (model 538) in a water bath (Julabo model F12) thermostated at a constant temperature maintained within ±0.01 °C. The electrolytic conductance was measured using a cell consisting of two platinum electrodes to which an alternating potential was applied. A conductometric cell with a cell constant of 0.84 cm⁻¹ was used throughout the studies.

3 Data analysis program

The 1:1 complexation of a metal cation, Mⁿ⁺, with a macrobicyclic ligand [L] is represented by the following equilibrium:

Here Λ_M is the molar conductance of the metal ion before addition of the ligand, Λ_ML is the molar conductance of the complexed ion, Λ_obs is the molar conductance of the solution during titration, C_L is the analytical concentration of the L added, and C_M is the analytical concentration of the metal ion. The complex formation constant, K_f, and the molar conductance of the complex, Λ_ML, were obtained by computer fitting of Eqs. (3) and (4) to molar conductance as a function of the ligand/metal cation mole ratio data, using a non-linear least-squares program GENPLOT (Genplot, 1989). All calculated stability constants are summarized in Table 1. The details of the calculation of the stability constants of the metal ion complexes by the conductometric method have been described elsewhere (Rounaghi et al., 1997).

4 Results and discussion

The complex formation between C₁₈H₃₆N₂O₆ and the ${\text{UO}}_2^{2+}$ cation in AN–DMF, AN–THF and AN–DCM binary mixtures was investigated by molar conductance changes upon addition of the ligand to the ${\text{UO}}_2^{2+}$ cation solution at 15, 25, 35 and 45 °C. Some of the resulting molar conductance (Λ_M) versus macrocycle to ${\text{UO}}_2^{2+}$ molar ratio ([L]_t/[M]_t) plots are shown in Figs. 2 and 3. As is evident from these figures, addition of C₁₈H₃₆N₂O₆ to a solution of the ${\text{UO}}_2^{2+}$ cation in AN–DCM (mol% AN = 75) and in AN–THF (mol% AN = 75) binary mixtures at different temperatures results in an increase in molar conductivity with an increase in the ligand concentration which indicates that the (C₁₈H₃₆N₂O₆·UO₂)²⁺ complex is more mobile than the free solvated ${\text{UO}}_2^{2+}$ cation. Similar behavior was observed for the other solvent systems.

Close

Figure 2

Molar conductance–mole ratio plots for (C₁₈H₃₆N₂O₆·UO₂)²⁺ complex in AN–DCM binary solution (%mol AN = 75) at different temperatures (♦ – 15 °C, ● – 25 °C, ▴ – 35 °C, ■ – 45 °C).

Export to PPT

Close

Figure 3

Molar conductance–mole ratio plots for (C₁₈H₃₆N₂O₆·UO₂)²⁺ complex in AN–THF binary solutions (%mol AN = 75) at different temperatures (♦ – 15 °C, ● – 25 °C, ▴ – 35 °C, ■ – 45 °C).

Export to PPT

The slope of the corresponding molar conductivity versus ligand/cation mole ratio plots changes at the point where the ligand to cation mole ratio is about 1, which is evidence for the formation of a relatively stable 1:1 complex. The role of solvent in complex formation reactions is convenient to discuss in terms of relationships between the stability constants of the complexes and the compositions of the mixed solvents (Rounaghi et al., 2002). For example, in the case of an aqueous non-aqueous binary mixed solvent solution, we can use the following equations:

Here K_exp[ML]ⁿ⁺ is the experimental stability constant of the complex in the binary mixed solvent with mole fraction X_s of the non-aqueous solvent, K_ex[ML]ⁿ⁺ is the additive stability constant, K_H2O[ML]ⁿ⁺ is the stability constant of the complex in water, K_s,max[ML]ⁿ⁺ is the stability constant in the mixed solvent with the maximum mole fraction of the non-aqueous component, and X_s,max is the maximum mole fraction of the non-aqueous solvent in the mixture, where K_exp[ML]ⁿ⁺ can be determined.

The results obtained for the complexation process between the ${\text{UO}}_2^{2+}$ cation and C₁₈H₃₆N₂O₆ show that the stoichiometry of the complex formed between the ligand and the ${\text{UO}}_2^{2+}$ cation is effected by the nature of the solvent systems. As evident from Fig. 4, in the AN–DMF (mol% AN = 75) binary system, addition of C₁₈H₃₆N₂O₆ to a solution of the cation at different temperatures causes molar conductivity to initially decrease until the molar ratio reaches 1:1 and then to increase. Similar behavior was observed in the case of AN–DMF binary mixed solvent with 50 mol% of AN. Such behavior may be described according to the following equilibria:

Close

Figure 4

Molar conductance–mole ratio plots for complex formation between C₁₈H₃₆N₂O₆ and ${\text{UO}}_2^{2+}$ in AN–DMF binary solutions (%mol AN = 75) at different temperatures (♦ – 15 °C, ● – 25 °C, ▴ – 35 °C, ■ – 45 °C).

Export to PPT

It seems that the addition of the ligand to the cation solution results in the formation of a relatively stable 1:1 (M:L) complex (Eq. (7)), in which the mobility of the solvated complex is close to the mobility of the free solvated ${\text{UO}}_2^{2+}$ cation, then the addition of the second ligand to the 1:1 [M:L] complex, causes the formation of a 1:2 [M:L] complex with a sandwich structure (Eq. (8)), which is less solvated than the 1:1 complex in dimethylformamide solutions and, therefore, conductivity increases.

The log K_f values in Table 1 show that the stability constant of the (C₁₈H₃₆N₂O₆·UO₂)²⁺ complex is larger in acetonitrile than dimethylformamide. The solvation of the ligand and the metal cation is influenced by the donor ability and dielectric constant of the solvent. It is known that the donor ability and the dielectric constant of the solvent play an important role in complexation reactions (Gutmann, 1978; Rounaghi et al., 2002). In a solvent with a high solvating ability (high donor number) such as DMF (DN = 26.6), complex formation tends to be weak, since the solvent solvates the cation strongly and competes with the ligand for the metal cation, but in solvents with lower donicity such as acetonitrile (DN = 14.1), the relatively poorer solvating ability of this solvent leads to an increase in the stability constant. Therefore, there is actually an inverse relationship between the stability of the complex and the donor ability of these organic solvents.

The changes of log K_f of the (C₁₈H₃₆N₂O₆·UO₂)²⁺ complex versus the mole fraction of THF in the AN–THF binary system at different temperatures are shown in Fig. 5. As shown in this figure, the change of the stability constant (log K_f) of the (C₁₈H₃₆N₂O₆·UO₂)²⁺ complex with the composition of AN–THF binary solution, is not linear. Somewhat similar behavior was also observed in the other binary solvent solutions. The non-monotonic behavior which is observed between log K_f and the composition of the mixed solvent solutions is due to solvent–solvent interactions between the constituent solvents which form the mixed solvent systems. The interaction between the solvent molecules in some binary mixed solvents has been studied (Krestov et al., 1994). For example, mixing of acetonitrile with dimethylformamide induces the mutual destruction of dipolar structures of these dipolar aprotic solvents and the release of the free dipoles (Rounaghi et al., 2007). As a result, a strong dipolar interaction between acetonitrile and dimethylformamide molecules is expected. In addition to solvent–solvent interactions, the preferential solvation of the cation, ligand and complex and also heteroselective solvation of these species in the binary mixed solvent solutions may be effective in non-linear relationship between log K_f and the composition of the mixed solvents. The effect of properties of AN–H₂O and other organic-water binary mixtures on the solvation enthalpy of 15C5 and some of the other crown ethers has been studied by Jozwaik (2004), and it has been shown that the crown ethers are preferentially solvated by organic solvents and it has been discussed that in these mixed solvents, the energetic effect of the preferential solvation depends quantitatively on the structural and energetic properties of the mixtures. Preferential solvation of ions by one of the components of a mixed solvent system depends on two factors: the relative donor–acceptor abilities of the component molecules toward the ion and the interactions between solvent molecules themselves. The solvating properties of the components in mixed solvents can even be significantly modified by solvent–solvent interactions when the energy of the latter is comparable with the energy difference of solvent–ion interactions for both components (Szymanska-Cybalska and Kamienska-Piotrowicz, 2006).

Close

Figure 5

Changes of stability constant (log K_f) of (C₁₈H₃₆N₂O₆·UO₂)²⁺ complex with the mole fraction of THF in AN–THF binary solvent solution at different temperatures (♦ = 15 °C, ■ = 25 °C, ▴ = 35 °C, × = 45 °C).

Export to PPT

Assuming that the activity coefficients of the cation and the complex have the same values, K_f, is a thermodynamic equilibrium constant on the molar concentration scale, related to the Gibbs energy of the complexation reaction, $\mathrm{\Delta }{G}_{\text{c}}^{°}$ . The van't Hoff plots of ln K_f versus 1/T in all cases were linear. The changes in the standard enthalpy ( $\mathrm{\Delta }{H}_{\text{c}}^{°}$ ) for the complexation of C₁₈H₃₆N₂O₆ with the ${\text{UO}}_2^{2+}$ cation were obtained from the slope of the van't Hoff plots assuming that ΔCp is equal to zero over the entire temperature range investigated. The changes in standard entropy ( $\mathrm{\Delta }{S}_{\text{c}}^{°}$ ) were calculated from the relationship ( $\mathrm{\Delta }{G}_{\text{c,}298.15}^{°}=\mathrm{\Delta }{H}_{\text{c}}^{°}-298.15\mathrm{\Delta }{S}_{\text{c}}^{°}$ ). The experimental values of standard enthalpy ( $\mathrm{\Delta }{H}_{\text{c}}^{°}$ ) and standard entropy ( $\mathrm{\Delta }{S}_{\text{c}}^{°}$ ) which are given in Table 2, show that in most cases, the complex is enthalpy stabilized and in all cases entropy stabilized.

As expected, the values of $\mathrm{\Delta }{H}_{\text{c}}^{°}$ and $\mathrm{\Delta }{S}_{\text{c}}^{°}$ depend strongly on the nature and composition of the mixed solvents. The value and the sign of the standard entropy changes are expected to vary with different parameters, such as changes in flexibility of the macrobicyclic ligand during the complexation process and the extent of cation–solvent, ligand–solvent and also the complex–solvent interactions. As is evident from Table 2, the standard thermodynamic parameters ( $\mathrm{\Delta }{H}_{\text{c}}^{°}$ and $\mathrm{\Delta }{S}_{\text{c}}^{°}$ ) change non-monotonically with the composition of the mixed solvents. Since there are many parameters which contribute to the changes of enthalpy and entropy of the complextion reactions, we should not expect to observe a monotonic behavior between these thermodynamic quantities and the solvent composition of these binary solvent solutions.